Acids Practice Questions

Hydrogen chloride gas dissolves in water to form hydrochloric acid. Name the ion that makes this solution acidic and give its formula.

Sodium hydroxide solution is alkaline. Name the ion that makes this solution alkaline.

State the pH of a neutral solution.

A solution of potassium hydroxide has a pH of 11. A different solution of sodium hydroxide has a pH of 13.

State which solution contains the higher concentration of hydroxide ions.

Complete the table to show the four missing colours.

The student adds phenolphthalein to an unknown solution and it stays colourless. Explain why the student cannot conclude that the unknown solution is acidic.

Name the gas produced in this reaction.

Describe the chemical test, and the positive result, that would confirm the identity of this gas.

The student then adds a piece of magnesium ribbon to dilute hydrochloric acid in a test tube. A different gas is produced.

Name this gas and state what would be seen when a lighted splint is held at the mouth of the test tube.

Use the solubility rules to complete the table. Put one tick (✓) in each row.

Most carbonates are insoluble in water. Name one carbonate that is soluble in water.

Name the salt formed when copper(II) oxide reacts with dilute sulfuric acid.

Name the salt formed when zinc reacts with dilute hydrochloric acid.

Calcium carbonate reacts with dilute nitric acid. Name the three products of this reaction.

State what is meant by the term base.

State what the student used to measure the pH values to one decimal place.

Give one variable the student must keep the same if the investigation is repeated with calcium oxide instead of calcium hydroxide.

Describe how the pH changes as the calcium hydroxide is added. Use data from the table in your answer.

Explain why the pH does not change after $0.50\ \mathrm{g}$ of calcium hydroxide has been added.

Which statement describes a dilute solution of a strong acid?

• A a small amount of acid dissolved in a large volume of water, and the acid is fully dissociated into ions
• B a large amount of acid dissolved in a small volume of water, and the acid is fully dissociated into ions
• C a small amount of acid dissolved in a large volume of water, and the acid is only partly dissociated into ions
• D a large amount of acid dissolved in a small volume of water, and the acid is only partly dissociated into ions

Explain what is meant by a weak acid.

A solution of hydrochloric acid has a pH of 1.0. Water is added to this solution until the hydrogen ion concentration is 100 times lower. Give the pH of the diluted solution.

A bottle is labelled 'concentrated ethanoic acid'. Explain why this solution is still described as a weak acid.

Write the balanced symbol equation, including state symbols, for the reaction of copper(II) carbonate with dilute nitric acid.

The technician keeps adding copper(II) carbonate until some solid remains unreacted at the bottom of the beaker. Explain why an excess of copper(II) carbonate is added.

Describe how the technician obtains dry crystals of copper(II) nitrate from the mixture in the beaker.

State why the dilute sulfuric acid is warmed before the copper(II) oxide is added.

State how the student can tell that enough copper(II) oxide has been added.

After filtering, the blue solution is heated in an evaporating basin placed over a water bath rather than being heated directly with a Bunsen burner. Give two reasons for using a water bath.

Describe how the crystals should be dried.

Give one safety precaution the student should take, and a reason for it.

Give one reason why the mass of crystals obtained is often less than the maximum possible.

Explain why potassium nitrate cannot be prepared by adding an excess of potassium hydroxide to the acid and then filtering.

In a titration, name the piece of apparatus used to measure exactly $25.0\ \mathrm{cm^3}$ of potassium hydroxide solution into the conical flask, and the piece of apparatus used to add the dilute nitric acid.

Once the volume of nitric acid needed has been found, describe how a pure, dry sample of potassium nitrate is obtained.

Complete the table using the solubility rules.

Explain, using the solubility rules, why no precipitate forms when sodium carbonate solution and potassium nitrate solution are mixed.

Write the ionic equation, including state symbols, for the formation of the precipitate in row 1 of the table.

Write the balanced symbol equation, including state symbols, for the reaction between iron and dilute sulfuric acid.

The solution formed contains iron(II) ions, $\mathrm{Fe^{2+}}$. Write the ionic equation for the reaction of iron with the hydrogen ions in the acid.

Write the ionic equation for the neutralisation of dilute nitric acid by potassium hydroxide solution.

Use solutions A and B to explain the relationship between hydrogen ion concentration and pH.

Explain why solution C has a higher pH than solution A, even though the two solutions have the same concentration.

Solution D is diluted with water until its concentration is $0.0010\ \mathrm{mol\ dm^{-3}}$. Predict the pH of the diluted solution and explain your answer.

Using the solubility rules, name two solutions that the student could mix to make lead(II) chloride. Justify each choice.

The diagram shows the apparatus used to separate the precipitate from the mixture.

Name the pieces of apparatus labelled A and B.

After the mixture has been filtered, describe how the student obtains a pure, dry sample of lead(II) chloride, explaining each step.

Both compounds are bases, but only sodium hydroxide is an alkali. Explain why.

Explain why copper(II) sulfate crystals can be prepared by adding an excess of copper(II) hydroxide to dilute sulfuric acid, but sodium sulfate crystals cannot be prepared by adding an excess of sodium hydroxide solution to dilute sulfuric acid.